Calcium sulfate

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Calcium sulphate
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Calcium sulfate
Calcium sulphate anhydrous
Calcium sulfate hemihydrate
Names
Other names
Sulfate of lime
Drierite
Gypsum
Identifiers
3D model (
JSmol
)
ChEBI
ChEMBL
ChemSpider
DrugBank
ECHA InfoCard
 100.029.000 Edit this at Wikidata
EC Number
  • 231-900-3
E number E516 (acidity regulators, ...)
7487
KEGG
RTECS number
  • WS6920000
  • (dihydrate): MG2360000
UNII
  • InChI=1S/Ca.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2 checkY
    Key: OSGAYBCDTDRGGQ-UHFFFAOYSA-L checkY
  • InChI=1/Ca.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
    Key: OSGAYBCDTDRGGQ-NUQVWONBAU
  • [Ca+2].[O-]S([O-])(=O)=O
Properties
CaSO4
Molar mass 136.14 g/mol (anhydrous)
145.15 g/mol (hemihydrate)
172.172 g/mol (dihydrate)
Appearance white solid
Odor odorless
Density 2.96 g/cm3 (anhydrous)
2.32 g/cm3 (dihydrate)
Melting point 1,460 °C (2,660 °F; 1,730 K) (anhydrous)
0.26 g/100ml at 25 °C (dihydrate)[1]
4.93 × 10−5 mol2L−2 (anhydrous)
3.14 × 10−5 (dihydrate)
[2]
Solubility in glycerol slightly soluble (dihydrate)
Acidity (pKa) 10.4 (anhydrous)
7.3 (dihydrate)
-49.7·10−6 cm3/mol
Structure
orthorhombic
Thermochemistry
107 J·mol−1·K−1 [3]
Std enthalpy of
formation
(ΔfH⦵298)
 -1433 kJ/mol[3]
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
1
0
0
Flash point Non-flammable
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 15 mg/m3 (total) TWA 5 mg/m3 (resp) [for anhydrous form only][4]
REL (Recommended)
 TWA 10 mg/m3 (total) TWA 5 mg/m3 (resp) [anhydrous only][4]
IDLH
(Immediate danger)
 N.D.[4]
Safety data sheet (SDS) ICSC 1589
Related compounds
Other cations
 Magnesium sulfate
Strontium sulfate
Barium sulfate
Related desiccants
 Calcium chloride
Magnesium sulfate
Related compounds
Plaster of Paris
Gypsum
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Calcium sulfate (or calcium sulphate) is the inorganic compound with the formula CaSO4 and related

permanent hardness
in water.

Hydration states and crystallographic structures

The compound exists in three levels of hydration corresponding to different crystallographic structures and to minerals:

  • CaSO
    4     (
    zirconium orthosilicate
    (zircon): Ca2+
     is 8-coordinate, SO2−
    4     is tetrahedral, O is 3-coordinate.
  • CaSO
    4    ·2H
    2    O     (gypsum and selenite (mineral)): dihydrate.[7]
  • CaSO
    4    ·1/2H
    2    O     (
    plaster of Paris. Specific hemihydrates are sometimes distinguished: α-hemihydrate and β-hemihydrate.[8]

Uses

See also: Gypsum § Uses

The main use of calcium sulfate is to produce plaster of Paris and stucco. These applications exploit the fact that calcium sulfate which has been powdered and calcined forms a moldable paste upon hydration and hardens as crystalline calcium sulfate dihydrate. It is also convenient that calcium sulfate is poorly soluble in water and does not readily dissolve in contact with water after its solidification.

Hydration and dehydration reactions

With judicious heating, gypsum converts to the partially dehydrated mineral called

plaster of Paris. This material has the formula CaSO4·(nH2O), where 0.5 ≤ n ≤ 0.8.[8]
Temperatures between 100 and 150 °C (212–302 °F) are required to drive off the water within its structure. The details of the temperature and time depend on ambient humidity. Temperatures as high as 170 °C (338 °F) are used in industrial calcination, but at these temperatures γ-anhydrite begins to form. The heat energy delivered to the gypsum at this time (the heat of hydration) tends to go into driving off water (as water vapor) rather than increasing the temperature of the mineral, which rises slowly until the water is gone, then increases more rapidly. The equation for the partial dehydration is:

CaSO4 · 2 H2O   →   CaSO4 · 1/2 H2O + 1+1/2 H2O↑

The

endothermic property of this reaction is relevant to the performance of drywall, conferring fire resistance to residential and other structures. In a fire, the structure behind a sheet of drywall will remain relatively cool as water is lost from the gypsum, thus preventing (or substantially retarding) damage to the framing (through combustion of wood members or loss of strength of steel at high temperatures) and consequent structural collapse. But at higher temperatures, calcium sulfate will release oxygen and act as an oxidizing agent. This property is used in aluminothermy
. In contrast to most minerals, which when rehydrated simply form liquid or semi-liquid pastes, or remain powdery, calcined gypsum has an unusual property: when mixed with water at normal (ambient) temperatures, it quickly reverts chemically to the preferred dihydrate form, while physically "setting" to form a rigid and relatively strong gypsum crystal lattice:

CaSO4 · 1/2 H2O + 1+1/2 H2O   →   CaSO4 · 2 H2O

This reaction is

cast earth, an alternative to adobe
(which loses its strength when wet). The conditions of dehydration can be changed to adjust the porosity of the hemihydrate, resulting in the so-called α- and β-hemihydrates (which are more or less chemically identical).

On heating to 180 °C (356 °F), the nearly water-free form, called γ-anhydrite (CaSO4·nH2O where n = 0 to 0.05) is produced. γ-Anhydrite reacts slowly with water to return to the dihydrate state, a property exploited in some commercial desiccants. On heating above 250 °C, the completely anhydrous form called β-anhydrite or "natural" anhydrite is formed. Natural anhydrite does not react with water, even over geological timescales, unless very finely ground.

The variable composition of the hemihydrate and γ-anhydrite, and their easy inter-conversion, is due to their nearly identical crystal structures containing "channels" that can accommodate variable amounts of water, or other small molecules such as methanol.

Food industry

The calcium sulfate hydrates are used as a coagulant in products such as tofu.[9]

For the

FDA, it is permitted in cheese and related cheese products; cereal flours; bakery products; frozen desserts; artificial sweeteners for jelly & preserves; condiment vegetables; and condiment tomatoes and some candies.[10]

It is known in the E number series as E516, and the UN's FAO knows it as a firming agent, a flour treatment agent, a sequestrant, and a leavening agent.[10]

Dentistry

Calcium sulfate has a long history of use in dentistry.[11] It has been used in bone regeneration as a graft material and graft binder (or extender) and as a barrier in guided bone tissue regeneration. It is a biocompatible material and is completely resorbed following implantation.[12] It does not evoke a significant host response and creates a calcium-rich milieu in the area of implantation.[13]

Other uses

Drierite

When sold at the anhydrous state as a desiccant with a color-indicating agent under the name

Drierite, it appears blue (anhydrous) or pink (hydrated) due to impregnation with cobalt(II) chloride
, which functions as a moisture indicator.

Up to the 1970s, commercial quantities of

clinker production.[14]

2 CaSO4 + 2 SiO2 + C → 2 CaSiO3 + 2 SO2 + CO2[15]

The plants made sulfuric acid by the “Anhydrite Process”, in which

catalyst
.

CaSO4 + 2 C → CaS + 2CO2

3 CaSO4 + CaS + 2 SiO2 → 2 Ca2SiO4 (belite) + 4 SO2

3 CaSO4 + CaS → 4 CaO + 4 SO2

Ca2SiO4 + CaO → Ca3OSiO4 (alite)

2 SO2 + O2 → 2 SO3 (in the presence of the

vanadium pentoxide
)

SO3 + H2O → H2SO4 [14]

Because of its use in an expanding niche market, the plant at Whitehaven, Cumbria continued to expand in a manner not shared by the other Anhydrite Process plants. The anhydrite mine opened on 11/1/1955, and the acid plant started on 14/11/1955. For a while in the early 1970s, it became the largest sulfuric acid plant in the UK, making about 13% of national production, and it was by far the largest Anhydrite Process plant ever built.[16]

Production and occurrence

The main sources of calcium sulfate are naturally occurring gypsum and anhydrite, which occur at many locations worldwide as evaporites. These may be extracted by open-cast quarrying or by deep mining. World production of natural gypsum is around 127 million tonnes per annum.[17]

In addition to natural sources, calcium sulfate is produced as a by-product in a number of processes:

  • In
    fossil-fuel power stations and other processes (e.g. cement manufacture) are scrubbed to reduce their sulfur dioxide content, by injecting finely ground limestone:[18]
SO2 + 0.5 O2 + CaCO3 → CaSO4 + CO2

Related sulfur-trapping methods use lime and some produces an impure calcium sulfite, which oxidizes on storage to calcium sulfate.

  • In the production of phosphoric acid from phosphate rock, calcium phosphate is treated with sulfuric acid and calcium sulfate precipitates. The product, called phosphogypsum is often contaminated with impurities making its use uneconomic.
  • In the production of hydrogen fluoride, calcium fluoride is treated with sulfuric acid, precipitating calcium sulfate.
  • In the refining of zinc, solutions of zinc sulfate are treated with hydrated lime to co-precipitate heavy metals such as barium.
  • Calcium sulfate can also be recovered and re-used from scrap drywall at construction sites.

These precipitation processes tend to concentrate radioactive elements in the calcium sulfate product. This issue is particular with the phosphate by-product, since phosphate ores naturally contain

lead-210 and polonium-210. Extraction of uranium from phosphorus ores can be economical on its own depending on prices on the uranium market or the separation of uranium can be mandated by environmental legislation and its sale is used to recover part of the cost of the process.[19][20][21]

Calcium sulfate is also a common component of fouling deposits in industrial heat exchangers, because its solubility decreases with increasing temperature (see the specific section on the retrograde solubility).

Retrograde solubility

The dissolution of the different crystalline phases of calcium sulfate in water is

endothermic (i.e., the reaction consumes heat: increase in Enthalpy: ΔH > 0) and whose solubility increases with temperature. Another calcium compound, calcium hydroxide (Ca(OH)2, portlandite
) also exhibits a retrograde solubility for the same thermodynamic reason: because its dissolution reaction is also exothermic and releases heat. So, to dissolve the maximum amount of calcium sulfate or calcium hydroxide in water, it is necessary to cool the solution down close to its freezing point instead of increasing its temperature.

Temperature dependence of the solubility of calcium sulfate (3 phases) in pure water.

The retrograde solubility of calcium sulfate is also responsible for its precipitation in the hottest zone of heating systems and for its contribution to the formation of scale in boilers along with the precipitation of calcium carbonate whose solubility also decreases when CO2 degasses from hot water or can escape out of the system.

On planet Mars

2011 findings by the Opportunity rover on the planet Mars show a form of calcium sulfate in a vein on the surface. Images suggest the mineral is gypsum.[22]

See also

References

  1. S2CID 132916752
    .
  2. ^ D.R. Linde (ed.) "CRC Handbook of Chemistry and Physics", 83rd Edition, CRC Press, 2002
  3. ^
    ISBN 978-0-618-94690-7
    .
  4. ^ a b c NIOSH Pocket Guide to Chemical Hazards. "#0095". National Institute for Occupational Safety and Health (NIOSH).
  5. doi:10.1002/14356007.a04_555
  6. doi:10.1107/S0567740875007145
    .
  7. doi:10.1107/S0567740874004055
    .
  8. ^
    ISBN 0-12-683900-X
    , pp. 186-187.
  9. ^ "About tofu coagulant". www.soymilkmaker.com. Sanlinx Inc. 31 August 2015. Archived from the original on 14 March 2015. Retrieved 10 January 2008.
  10. ^ a b "Compound Summary for CID 24497 - Calcium Sulfate". PubChem.
  11. ISSN 0032-5791
    .
  12. PMID 16393128
    .
  13. ^ "Biphasic Calcium Sulfate - Overview". Augma Biomaterials. 2020-03-25. Retrieved 2020-07-16.
  14. ^ a b Anhydrite Process
  15. ^ COMMONWEALTH OF AUSTRALIA. DEPARTMENT OF SUPPLY AND SHIPPING. BUREAU OF MINERAL RESOURCES GEOLOGY AND GEOPHYSICS. REPORT NO.1949/44 (Geol. Ser. No. 27) by E.K. Sturmfels THE PRODUCTION OF SULPHURIC ACID AND PORTLAND CEMENT FROM CALCIUM SULPHATE AND ALUMINIUM SILICATES
  16. ^ Whitehaven Cement Plant
  17. ^ Gypsum, USGS, 2008
  18. ISBN 9780471484943
    .
  19. OSTI 6654998
    .
  20. ^ "Uranium from Phosphates | Phosphorite Uranium - World Nuclear Association".
  21. ^ "Brazil plans uranium-phosphate extraction plant in Santa Quitéria : Uranium & Fuel - World Nuclear News".
  22. ^ "NASA Mars Opportunity rover finds mineral vein deposited by water". NASA Jet Propulsion Laboratory. December 7, 2011. Retrieved April 23, 2013.

External links

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